Isotope

Isotopes are any of the several different forms of an element each having different atomic mass (mass number). Isotopes of an element have nuclei with the same number of protons (the same atomic number) but different numbers of neutrons. Therefore, isotopes have different mass numbers, which give the total number of nucleons—the number of protons plus neutrons.

A nuclide is any particular atomic nucleus with a specific atomic number Z and mass number A; it is equivalently an atomic nucleus with a specific number of protons and neutrons. Collectively, all the isotopes of all the elements form the set of nuclides. The distinction between the terms isotope and nuclide has somewhat blurred, and they are often used interchangeably. Isotope is best used when referring to several different nuclides of the same element; nuclide is more generic and is used when referencing only one nucleus or several nuclei of different elements. For example, it is more correct to say that an element such as fluorine consists of one stable nuclide rather than that it has one stable isotope.

In IUPAC nomenclature, isotopes and nuclides are specified by the name of the particular element, implicitly giving the atomic number, followed by a hyphen and the mass number (e.g. helium-3, carbon-12, carbon-13, iodine-131 and uranium-238). In symbolic form, the number of nucleons is denoted as a superscripted prefix to the chemical symbol (e.g. ³He, ¹²C, ¹³C, ¹³¹I and ²³⁸U).

The term isotope was coined in 1913 by Margaret Todd, a Scottish doctor, during a conversation with Frederick Soddy (to whom she was distantly related by marriage).^[1] Soddy, a chemist at Glasgow University, explained that it appeared from his investigations as if several elements occupied each position in the periodic table. Hence Todd suggested the Greek for "at the same place" as a suitable name. Soddy adopted the term and went on to win the Nobel Prize for Chemistry in 1921 for his work on radioactive substances.

In the bottom right corner of JJ Thomson's photographic plate are markings for the two isotopes of neon: neon-20 and neon-22.

In 1913, as part of his exploration into the composition of canal rays, JJ Thomson channeled a stream of ionized neon through a magnetic and an electric field and measured its deflection by placing a photographic plate in its path. Thomson observed two patches of light on the photographic plate (see image on right), which suggested two different parabolas of deflection. Thomson concluded that some of the atoms in the gas were of higher mass than the rest.

Variation in properties between isotopes

A neutral atom has the same number of electrons as protons. Thus, different isotopes of a given element all have the same number of protons and electrons and the same electronic structure; because the chemical behavior of an atom is largely determined by its electronic structure, isotopes exhibit nearly identical chemical behavior. The main exception to this is the kinetic isotope effect: due to their larger masses, heavier isotopes tend to react somewhat more slowly than lighter isotopes of the same element.

This "mass effect" is most pronounced for protium (¹H) vis-à-vis deuterium (²H), because deuterium has twice the mass of protium. For heavier elements the relative mass difference between isotopes is much less, and the mass effect is usually negligible.

Similarly, two molecules which differ only in the isotopic nature of their atoms (isotopologues) will have identical electronic structure and therefore almost indistinguishable physical and chemical properties (again with deuterium providing the primary exception to this rule). The vibrational modes of a molecule are determined by its shape and by the masses of its constituent atoms. Consequently, isotopologues will have different sets of vibrational modes. Since vibrational modes allow a molecule to absorb photons of corresponding energies, isotopologues have different optical properties in the infrared range.

Although isotopes exhibit nearly identical electronic and chemical behavior, their nuclear behavior varies dramatically. Atomic nuclei consist of protons and neutrons bound together by the strong nuclear force. Because protons are positively charged, they repel each other. Neutrons, which are electrically neutral, allow some separation between the positively charged protons, reducing the electrostatic repulsion. Neutrons also stabilize the nucleus because at short ranges they attract each other and protons equally by the strong nuclear force, and this also offsets the electrical repulsion between protons. For this reason, one or more neutrons are necessary for two or more protons to be bound into a nucleus. As the number of protons increases, additional neutrons are needed to form a stable nucleus; for example, although the neutron to proton ratio of ³He is 1:2, the neutron/proton ratio of ²³⁸U is greater than 3:2. If too many or too few neutrons are present, the nucleus is unstable and subject to nuclear decay.ISOTOPES ARE AWSOME' link title

Occurrence in nature

Most elements have several different isotopes that can be found in nature. The relative abundance of an isotope is strongly correlated with its tendency toward nuclear decay; short-lived nuclides quickly decay away, while their long-lived counterparts endure. However, this does not mean that short-lived species disappear entirely; many are continually produced through the decay of longer-lived nuclides. Also, short-lived isotopes such as those of promethium have been detected in the spectra of stars, where they presumably are being continuously made by stellar nucleosynthesis. The tabulated atomic masses of elements are averages that account for the presence of multiple isotopes with different masses.

According to generally accepted physical cosmology|cosmology, virtually all nuclides other than isotopes of hydrogen and helium (and traces of some isotopes of lithium, beryllium and boron-- see big bang nucleosynthesis) were built in stars and supernovae. Their respective abundances here result from the quantities formed by these processes, their spread through the galaxy, and their rates of decay. After the initial coalescence of the solar system, isotopes were redistributed according to mass. The isotopic composition of elements is different on different planets, making it possible to determine the origin of meteorites.

Molecular mass of isotopes

The molecular mass (M_r) of an element is determined by its nucleons. For example, Carbon-12 (¹²C) has 6 Protons and 6 Neutrons. When a sample contains two isotopes the equation below is applied where M_r(1) and M_r(2) are the molecular masses of each individual isotope, and % abundance is the percentage abundance of that isotope in the sample.

M_{r}={\frac {M_{r}(1)*\%abundance+M_{r}(2)*\%abundance}{100}}.

Applications of isotopes

Several applications exist that capitalize on properties of the various isotopes of a given element.

Use of chemical properties

One of the most common applications is isotopic labeling, the use of unusual isotopes as tracers or markers in chemical reactions. Normally, atoms of a given element are indistinguishable from each other. However, by using isotopes of different masses, they can be distinguished by mass spectrometry or infrared spectroscopy (see "Properties"). For example, in 'stable isotope labeling with amino acids in cell culture (SILAC)' stable isotopes are used to quantify proteins. If radioactive isotopes are used, they can be detected by the radiation they emit (this is called radioisotopic labeling).

A technique similar to radioisotopic labelling is radiometric dating: using the known half-life of an unstable element, one can calculate the amount of time that has elapsed since a known level of isotope existed. The most widely known example is radiocarbon dating used to determine the age of carbonaceous materials.

Isotopic substitution can be used to determine the mechanism of a reaction via the kinetic isotope effect.

Use of nuclear properties

Several forms of spectroscopy rely on the unique nuclear properties of specific isotopes. For example, nuclear magnetic resonance (NMR) spectroscopy can be used only for isotopes with a nonzero nuclear spin. The most common isotopes used with NMR spectroscopy are ¹H, ²D,¹⁵N, ¹³C, and ³¹P.

Mössbauer spectroscopy also relies on the nuclear transitions of specific isotopes, such as ⁵⁷Fe.

Radionuclides also have important uses. Nuclear power and nuclear weapons development require relatively large quantities of specific isotopes. The process of isotope separation represents a significant technological challenge, but more so with heavy elements such as uranium or plutonium, than with lighter elements such as hydrogen, lithium, carbon, nitrogen, and oxygen. The lighter elements are commonly separated by gas diffusion of their compounds such as CO and NO. Uranium isotopes have been separated in bulk by gas diffusion, gas centrifugation, laser ionization separation, and (in the Manhattan Project) by a type of production mass spectroscopy.

References

^ Budzikiewicz H, Grigsby RD (2006). "Mass spectrometry and isotopes: a century of research and discussion". Mass spectrometry reviews. 25 (1): 146–57. doi:10.1002/mas.20061. PMID 16134128.

External links

[pmid16134128-1] Budzikiewicz H, Grigsby RD (2006). "Mass spectrometry and isotopes: a century of research and discussion". Mass spectrometry reviews. 25 (1): 146–57. doi:10.1002/mas.20061. PMID 16134128.

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